Bonding

• Atoms are unstable due to their outer e- = valence e-.

• Remainder of the atom = kernel of the atom.

• Atoms react to become stable.

• Atoms can react to become stable in a number of ways = bonding

Types of Bonding:

I) Metallic Bonding

Elements involved =soft metals (all metals to the left of the steps except the center of d and the bottom of

p).

Characteristics of metals

1) Low electronegativity

• Little attraction for shared e-.

2) Low ionization en (IE)

• Little en needed to rip off an e-.

3) Low electron affinity (EA)

• Little desire to gain e-

4) Big

• Not holding e- very tightly.

5) Few reactive e-

• If lost will be stable

• Because of all the characteristics of metals they like to lose e- in order to become stable.

• In metallic bonding all metals want to lose their valence e-, but there is no atom to take them.

• So they promote their valence e- into en levels greater than the 7th.

• The host atom (the atom promoting its e-) then feels like it lost these e- (even though they didn’t

lose them) forming a cation (+).

• The whole thing sticks together because of the attraction between the (+) ions and the (-) e- from neighboring atoms forming a metallic

bond!

• Lattice (an orderly arrangement of particles)

• Draw and labelDraw and label

• Results and characteristics for metallic bonds:

• 1) Moderate in strength

• 2) Malleable

• 3) Ductile

• 4) Moderate m.p. and b.p.

• 5) Good conductor of heat

(They are fairly closely packed and energy is generated due to friction and is easily transferred

due to the tightly packed particles).

• 6) Good conductors of electricity

•Mobile e- can flow, experience a collision, and the last e- flies off.

• 7) Luster

(II) Pure (nonpolar) covalent bonds:

• Elements involved: nonmetals

• Located in the upper right hand side of the p. steps.

• Electronegativity difference

• Covalent (<1.67) = sharing

• Pure (nonpolar) = zero = equal.

•Properties of nonmetals:

1) high electronegativity

• strong attraction to share e-.

•2) high IE

•Difficult to rip off e-

3) high EA

• Strong desire to gain e-

4) small

• Holding on to their e- tightly

5) 5 or more valence e-

• Easier to gain than lose e- to become stable.

•Nonmetals want to keep their own e- and gain more!

•Do examples from notes 1) HDo examples from notes 1) H22 2) F 2) F22 3) O 3) O22 4) N 4) N22!!

•4 Types of Overlaps:

1) s + s

(ex) H2

2) s + p

(ex) HCl

3) p + p (head on)

(ex) F2

4) p + p (not head on)

(ex) O2

3) Polar Covalent Bonds

• 2 elements with different electronegativities

• Covalent - electronegativitiy difference (< 1.67)

• Polar - not equal to zero.

•A group of atoms (composed of nonmetals) that is united by covalent bonds is called a molecule and a

substance that is made of molecules is called a molecular substance.

•Do examples on the board 1) HCl and 2) NHDo examples on the board 1) HCl and 2) NH33

Properties of Covalent Bonds

- low m.p. & low b.p-many gases, solids or liquids at

room T (made of molecules) - poor electrical conductors in all

states of matter – many soluable in nonpolar liquids

but not in water

4) Ionic bonds

• Elements involved = metal and nonmetal

• Electronegativity diff. >1.67

• Most ideal type of bond.

• Each element is getting exactly what they want: metal is losing e- and the nonmetal is gaining e-.

Do examples on the board 1) NaCl 2) CaFDo examples on the board 1) NaCl 2) CaF22

3) AlCl3) AlCl33

Properties of elements involved in ionic bonding:

1) high m.p. and b.p

2) good electrolyte ( a substance whose water solution conducts an electric current) when

melted or dissolved because it breaks down into ions (charges). Noncoductors as solids.

3) Crystalline solids (made of ions)

4) Many soluable in water, but not in nonpolar liquids.

Comparison of Properties:

• Ionic compounds vs

• 1) crystalline solids (made of ions)

• 2) high mp and bp

• 3) conduct electricity when melted

• 4) many soluable in water but not in nonpolar liquids

Covalent comounds1) gases, lqs, and

solids (made of molecules)

2) low mp and bp3) Poor electrical

conductors in all phases

4) Many soluble in nonpolar liquids but not in water

• The types of attractions that exist between molecules depend upon the types of bonds

between the atoms of each molecule.

• The strength of these attractive forces determine the physical properties of the

compound, such as its melting point, hardness, and solubility.

• Both the percent ionic character of the bonds and the molecular geometry determine the overall

polarity of the molecule.

•For nonpolar covalent molecules, the only forces of attraction between the molecules is called

dispersion interaction.

• A dispersion interaction force is a very weak attraction between the positive nuclei of one molecule for the valence e- of a neighboring

molecule. A temporary imbalance of e- distribution occurs in the neighboring molecule

resulting in an induced dipole.

• CH4 is a nonpolar covalent molecule so the only forces of attraction that exists in this molecule

is dispersion interaction forces.

• Since this force is dependent upon the number of protons and electrons that are attracting each other, molecules with larger molar mass will have the

greatest dispersion interaction.

(ex) C8H18 vs. C50H102

• Polar molecules arrange themselves so that the opposite charges of adjacent dipoles are aligned

with each other. The attraction of the partial positive charge of one dipole for the partial negative charge of another dipole is called a

dipole-dipole attraction.

• Molecules, such as HF, that have a positively charged end and a negatively charged end are

referred to as dipoles.

• Van der Waals forces = dispersion interaction forces and dipole-dipole forces.

• Polar molecules that contain hydrogen atoms bonded to a highly electronegative element with

extra pairs of unshared valence e- exhibit additional attractive forces called hydrogen

bonds. (ex) HF, H2O

• Strongest to weakest bonding:

•Ionic bonding or metallic bonding

•hydrogen bonding

•Dipole - dipole attractions

•Dispersion interaction

• Attractive forces between molecules are the strongest as the molecules get closer together.

• Bonds between molecules are called intermolecular forces.

• Bonds within molecules are called intramolecular forces.

• Summary

•1) Pure covalent

• Only intermolecular force is dispersion interaction (very weak force between nuclei and valence e- of

neighboring molecules).

• Strength of the dispersion interaction force increases with an increase in formula mass.

• 2) Polar covalent

• Dipole-dipole attraction

(partial (+) and partial (-) charges line up),

• Also has dispersion interaction forces

• Van der waals forces - both dipole-dipole and dispersion interaction forces.

• Can have H bonding (H bonded to a highly electronegative element - F,N,O)

Oxidation/Reduction (Oil Rig)

• (oil = oxidation is a loss of e-)

• (rig = reduction is a gain of e-).

• (ex) Na + en -----> 1e- + Na1+

• (Na lost e-, therefore underwent oxidation)

• (ex) Cl + 1e- -----> en + Cl1-

• (Cl gained e-, therefore it underwent reduction)

• Na provided the e- to Cl and caused the reduction and is therefore the reducing agent.

• Cl took the e- from Na and caused the oxidation and is therefore the oxidizing agent).

Pair Repulsion

unshared-unshared > unshared-shared > shared-

shared

Use HUse H22O as an exampleO as an example

• Sometimes we see strange bonding that needs to be explained and we use hybridization to explain.

• Hybridization - the combining of 2 or more orbitals of nearly the same en into orbitals of exactly the same

en.

• Atoms do this by promoting e- from a pair into the closest empty orbital.

• Shape - if there are only 2 atoms the molecule’s shape is linear!

•1) BeF2

• Shape = linear (180)

• Hybrid = sp

• To determine the hybrid - count only lone pairs and sigma bonds.

• To determine the shape - look at the arrangement of the atoms directly off the main atom.

• 2) H2O

• Shape = bent (depends)

• Hybrid = none

• 3) BF3

• Shape = trigonal planar (120) - three atoms off the central atom

• hybrid = sp2

• 4) NH3

• Shape = trigonal pyramidal (<120) - three atoms off the central atom and an unshared pair of e-.

• hybrid = none

• 5) CH4

• Shape = tetrahedral (a 4 sided figure of which each side is an equalaterial triangle). (109.5) - 4 atoms off

the central atom.

•hyrid = sp3

6) PCl5

Shape = trigonal bipyramidal - 5 atoms off the central atom.

hybrid = sp3d

7) SF6

Shape = octahedral - 6 atoms off the central atom.

hybrid =sp3d2

(exs) C2H4 (1 double bond)

C2H2 (1 triple bond)

C2H6 (all single bonds)

• Molecules with the following geometric shapes:

• 1) linear (if the same element = O2)

• 2) trigonal planar (if atoms off the central element are the same = BF3).

• 3) tetrahedral (if atoms off the central atom are the same = CH4)

• Are nonpolar!

• Molecules with the remaining geometric shapes are polar!

• (exs) SeCl4, PH3, SiO2, O2, N2, HF

• Isomers - compounds that have the same molecular formula but different structural

formulas.

1) Structual isomers - C chain is altered.

(ex) C-C-C-C

C-C-C

|

C

2) Geometric isomers - different arrangement around a double bond.

(ex) C2H2Cl2 (cis and trans)

3) Positional isomers - diff position on the C chain.

(ex) C-C-C-OH

C-C-C

|

OH

4) Functional isomers - new element bonded in a diff way.

(ex) C-C-C-OH

C-C-O

|

C